The change in enthalpy of a chemical reaction indicates how much heat is absorbed or released by the system. This is valuable information in chemistry, because the exchange in heat affects the reaction conditions and the surroundings, and that needs to be managed and taken into account – in theory, in the laboratory, in industry or in nature in general.

Chemists often want to compare the changes in enthalpy between different reactions. Since changes in enthalpy depend on both temperature and pressure, we need to control for these 2 confounding variables by using a reference set of temperature and pressure. This set of conditions is called the standard conditions, and it sets the standard temperature at 298 degrees Kelvin and the standard pressure at 1 bar. (IUPAC changed the definition of standard pressure from 1 atmosphere to 1 bar in 1982. The actual difference in pressure between these 2 definitions is very small.)

The standard enthalpy of reaction (or standard heat of reaction) is the change in enthalpy of a chemical reaction under standard conditions; the actual number of moles are specified by the coefficients of the balanced chemical equation. (Since enthalpy is an extensive property, the same reaction under standard conditions could have different changes in enthalpy with different amounts of the reactants and products. Thus, the number of moles of the reaction must be standardized somehow when defining the standard enthalpy of reaction.) The standard enthalpy of reaction has the symbol ΔHº; the º symbol indicates the standard conditions.

Let’s examine the difference between a change enthalpy and a change in internal energy. It helps to think of the following 2 scenarios.

If the chemical reaction releases a gas but occurs at constant volume, then there is no pressure-volume work. The only way for energy to be transferred between the system and the surroundings is through heat. An example of a system under constant volume is a bomb calorimeter. In this case,

This heat is denoted as to indicate that this is heat transferred under constant volume. In this case, the change in enthalpy is the same as the change in internal energy.

If the chemical reaction releases a gas and occurs at constant pressure, then energy can be transferred between the system and the surroundings through heat and/or work. Thus,

This heat is denoted as to indicate that this is heat transferred under constant pressure. Thus, as the gas forms inside the cylinder, the piston pushes against the constant pressure that the atmosphere exerts on it. The total energy released by the chemical reaction allows some energy to be used for the pressure-volume work, with the remaining energy being released via heat. (Recall that these are the 2 ways for internal energy to be changed according to the First Law of Thermodynamics.) Thus, the difference between enthalpy and internal energy arises under constant pressure – the difference is the pressure-volume work.

Reactions under constant pressure are often illustrated by a reaction that releases a gas in cylinder with a movable piston, but they are actually quite common. In fact, in chemistry, reactions under constant pressure are much more common than reactions under constant volume. Chemical reactions often happen in beakers, flasks or any container open to the constant pressure of the atmosphere.

As the gas expands inside the cylinder, it pushes against the piston, and work is done by the system against the surroundings. The atmospheric pressure on the cylinder remains constant while the cylinder expands, and the volume of the cylinder increases as a result. The volume of the cylinder at any given point is the area of the piston times the length of the cylinder. The change in volume is equal to the area of the piston times the distance along which the piston was pushed by the expanding gas.

Note that this last line is just the definition of work under constant force in the same direction as the displacement, multiplied by the negative sign to follow the sign convention in chemistry.

Introduction

The Ideal Gas Law, , is a very simple yet useful relationship that describes the behaviours of many gases pretty well in many situations. It is “Ideal” because it makes some assumptions about gas particles that make the math and the physics easy to work with; in fact, the simplicity that arises from these assumptions allows the Ideal Gas Law to be easily derived from the kinetic theory of gases. However, there are situations in which those assumptions are not valid, and, hence, the Ideal Gas Law fails.

Boyle’s law is inherently a part of the Ideal Gas Law. It states that, at a given temperature, the pressure of an ideal gas is inversely proportional to its volume. Equivalently, it states the product of the pressure and the volume of an ideal gas is a constant at a given temperature.

An Example of The Failure of the Ideal Gas Law

This law is valid for many gases in many situations, but consider the following data on the pressure and volume of 1.000 g of oxygen at 0 degrees Celsius. I found this data set in Chapter 5.2 of “General Chemistry” by Darrell Ebbing and Steven Gammon.

The right-most column is the product of pressure and temperature, and it is not constant. However, are the differences between these values significant, or could it be due to some random variation (perhaps round-off error)?

Here is the scatter plot of the pressure-volume product with respect to pressure.

These points don’t look like they are on a horizontal line! Let’s analyze these data using normal linear least-squares regression in R.

In the early morning, especially here in Canada, I often see dew – water droplets formed by the condensation of water vapour on outside surfaces, like windows, car roofs, and leaves of trees. I also sometimes see fog – water droplets or ice crystals that are suspended in air and often blocking visibility at great distances. Have you ever wondered how they form? It turns out that partial pressure, vapour pressure and temperature are the key phenomena at work.